Pi bond

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Electron atomic and molecular orbitals, showing a Pi-bond at the bottom right of the picture.

The Greek letter π in their name refers to p orbitals, since the orbital symmetry of the pi bond is the same as that of the p orbital when seen down the bond axis. P orbitals usually engage in this sort of bonding. However, d orbitals can engage in pi bonding also.

Pi bonds are usually weaker than sigma bonds because their (negatively charged) electron density is farther from the positive charge of the atomic nucleus, which requires more energy. From the perspective of quantum mechanics, this bond's weakness is explained by significantly less overlap between the component p-orbitals due to their parallel orientation.

Although the pi bond by itself is weaker than a sigma bond, pi bonds are often components of multiple bonds, together with sigma bonds. The combination of pi and sigma bond is stronger than either bond by itself. The enhanced strength of a multiple bond vs. a single (sigma bond) is indicated in many ways, but most obviously by a contraction in bond lengths. For example in organic chemistry, carbon-carbon bond lengths are ethane (154 pm), ethylene (133 pm) and acetylene (120 pm).

Two s-orbitals continue to overlap when bond rotates because orientation is along axis. Circles represent s orbitals. Ellipses represent merged sigma bond. Pink and gray represent a ball and stick model of the molecular fragment that contains the sigma bond.

In addition to one sigma bond, a pair of atoms connected via double bond and triple bonds have one or two pi bonds, respectively. Pi bonds result from overlap of atomic orbitals that with two areas of overlap. Pi-bonds are more diffuse bonds than the sigma bonds. Electrons in pi bonds are sometimes referred to as pi electrons. Molecular fragments joined by a pi bond cannot rotate about that bond without breaking the pi bond, because rotation involves destroying the parallel orientation of the constituent p orbitals.

Special cases

Pi bonds do not necessarily connect a pair of atoms that are also sigma-bonded.

In certain metal complexes, pi interactions between a metal atom and alkyne and alkene pi antibonding orbitals form pi-bonds.

In some cases of multiple bonds between two atoms, there is no sigma bond at all, only pi bonds. Examples include diiron hexacarbonyl (Fe2(CO)6), dicarbon (C2) and the borane B2H2. In these compounds the central bond consists only of pi bonding, and in order to achieve maximum orbital overlap the bond distances are much shorter than expected.[1]